Understanding the Periodic Table of Elements

The Periodic Table

• Elements are arranged in groups based on similar chemical properties and reactivity
• Elements in a period share the same number of atomic shells or orbits

Periodic Table Labels

• Atomic number represents the number of protons (Z) in the nucleus.
• Atomic mass (A) is the average mass number

Atomic Components

• The nucleus contains protons (positive charge) and neutrons (no charge).
• The nucleus holds most of the atom's mass.
• Protons determine the element.
• Electrons are negatively charged and orbit the nucleus
• Electrons are much smaller than protons and neutrons.
• Electrons are attracted to the positive nucleus

Elements

• Elements consist of one type of atom
• Neutron number is calculated by subtracting the atomic number from the mass number: 95.94 - 42 = 53.94.

Families/Groups

• Alkali metals are in Group 1 and have one valence electron; they are shiny and soft.
• Alkaline earth metals are never found uncombined and have two valence electrons, like magnesium and calcium.
• Transition metals comprise 10 groups of metals between Groups 2 and 13, conducting heat and electricity.
• Halogens in the halogen family have seven valence electrons, are active non-metals, and form salts with alkali metals, such as sodium chloride (NaCl).
• Noble gases are un-reactive, colorless gases found in small amounts in the atmosphere, including helium, neon, argon, krypton, xenon, and radon.
• Rare earth elements include the lanthanide and actinide series.

Electron Configuration

• Electrons orbit the nucleus in electron shells or energy levels.
• The first shell has the strongest attraction and lowest energy.
• Electron configuration describes electron arrangement in shells.
• Shell capacity is given by 2n², where n is the shell number; shells fill sequentially
• Reactivity depends on achieving a full outer shell through valence electrons.

Isotopes

• Isotopes of an element have the same number of protons but different neutron numbers, leading to different atomic weights.
• Isotopes are named by atomic mass, e.g., carbon-12.
• Mass equals protons plus neutrons; atoms equal protons plus electrons

Radioisotopes

• Radioisotopes are unstable atoms that undergo nuclear decay by releasing particles or electromagnetic radiation.

Nuclear Radiation Types:

• Alpha (α) radiation involves emitting a cluster of 2 protons and 2 neutrons (helium nucleus) to increase stability.
• Beta (β) radiation occurs when a neutron changes into a proton and electron due to an imbalance.
• Gamma (γ) radiation is the emission of high-frequency electromagnetic radiation (e.g., X-rays) to stabilize the atom.

Chemical Reactions

• Chemical reactions rearrange atoms to create new substances.
• Chemical equations represent these reactions.
• Reactants are the starting substances.
• Products are the substances formed.
• Reactants are written on the left side and products on the right side of the arrow.

Reading Chemical Formulas

• Elements written together form a compound.
• Subscripts indicate the quantity of each element.

Chemical Formula Examples:

• NaHCO3: contains 1 sodium, 1 hydrogen, 1 carbon, and 3 oxygen atoms.
• Cu3(PO4)2: contains 3 copper, 2 phosphorus, and 8 oxygen atoms (multiply bracketed atoms by the subscript).
• CH3COOH: contains 2 carbon, 4 hydrogen, and 2 oxygen atoms; written separately to show structure.

Balancing Equations

• Chemical equations must be balanced to conserve mass.
• Add coefficients (numbers in front of chemicals) to balance the number of elements.

Reaction Types:

• Exothermic reactions release energy as heat or light.
• Endothermic reactions absorb energy from surroundings.
• Combination/Synthesis reactions form complex products from simple reactants
• Decomposition reactions break down complex reactants into simpler products.
• Single displacement involves one atom or group replacing another in a compound.
• Double displacement occurs when two compounds swap partners.
• Combustion is burning, an exothermic oxidation reaction involving a substance reacting with oxygen to form oxides, releasing energy.

Combustion Equations:

• Word equation: Magnesium + oxygen → Magnesium oxide.
• Balanced equation example: 2Mg + O2 → 2MgO.
• Substances containing only hydrogen and carbon atoms are hydrocarbons.
• Hydrocarbons include natural gas (CH4), petrol octane (C8H18), and candle wax (C20H42).

Combustion Equations Types

• Complete Combustion: Hydrocarbon + Oxygen → Carbon dioxide + water + energy, and requires sufficient oxygen
• Incomplete Combustion: occurs with limited oxygen, producing carbon monoxide and sooty carbon.

Acid + Metal Reactions

• Acid + Metal → Salt + Hydrogen gas
• The salt formed depends on the acid used

Acids

• Acids are chemicals containing hydrogen that dissolve in water to produce positive hydrogen ions (H+), exemplified by hydrochloric acid (HCl).

Bases (Alkali)

• Bases dissolve in water to form hydroxide ions (OH-), e.g., metal oxides and metal hydroxides like sodium hydroxide (NaOH).

pH Scale

• pH measures the concentration of hydrogen ions.

Neutralization

• Neutralization occurs when acids and bases react.
• Hydrogen ions (H+) from acids combine with hydroxide ions (OH-) from bases to form water (pH 7).
• Salts from neutralization are formed from leftover ions, where the acid provides the negative ion and the base the positive ion.

Respiration and Corrosion

• Aerobic respiration is a reaction producing energy in cells.
• Respiration and combustion both require oxygen and release carbon dioxide and energy and both require oxygen as a reactant, produce carbon dioxide as a product and release energy
• Respiration occurs slowly and does not release light.
• Combustion releases lots of heat and light
• Corrosion happens when metal oxides combine with oxygen to form metal oxides.
• Corrosion breaks metals down into compounds and releases heat slowly

Ions and Ionic Compounds

• Ions are atoms that have gained or lost electrons
• If the atom has more electrons than protons, it is a negative ion, or ANION, and tends to be non-metals(except hydrogen)
• If it has more protons than electrons, it is a positive ion (CATION), and tends to be metals.
• Ionic compounds form when atoms donate electrons to one another, such as sodium chloride.
• Cl wants to gain or receive an electron
• Ionic name: they form an ionic bond forming Sodium chloride – add ‘ide’ to the end of the anion

Polyatomic Ions

• Polyatomic ions are consist of 2 or more atoms with an overall change

Polyatomic ions examples:

• Hydroxide (OH)-
• Ammonium (NH4)+
• Nitrate (NO3)-
• Carbonate (CO3)2-
• Phosphate (PO4)3-
• Sulfate (SO4)2-
• Chemical equations of ionic compounds swap / 'cross over' the charges of the two ions
• If this involves a polyatomic ion, use brackets

Precipitation Reactions

• Precipitation reactions form an insoluble solid when two solutions are mixed
• Occurs when compounds 'swapping partners' i.e. double displacement reaction
• Reaction begins with 2 clear solutions – both solutions are soluble compounds
• When mixed, double displacement occurs. One of the new products must be insoluble (precipitate), hence appears to ‘fall out’ in the liquid
• Can predict what the precipitate will be by looking at a solubility table

Collision Theory and Reaction Rates

• Rate of reaction is the speed at which reaction takes place
• chemical reactions occur when different atoms or molecules collide, The particles must have:
• A certain amount of energy - this is called ACTIVATION ENERGY The correct orientation

Collision Theory

• Reactions occur when molecules collide with appropriate energy to break/form bonds.
• Four factors influence reaction rate:

Factors Influencing reaction rate:

• Temperature: increased temperature causes more kinetic energy and faster collisions.
• Concentration: increasing concentration means more particles and more collisions.
• Surface area: increasing surface area increases collisions.
• Catalyst: catalysts speed up reactions but are not consumed.
• Agitation (mixing or stirring) ensures the reactants are kept in contrast, by removing the build-up of products around the reactants